Chemistry. - Velocities of Nitration 1) (Provisional Communication). By F. H. COHEN. (Communicated by Prof. A. F. HOLLEMAN).

(Communicated at the meeting of June 30, 1928).

I. Introduction.

Only very few kinetic measurements of nitrations have been made as yet. GIERSBACH and KESSLER 2) and BRÄUER 3) have studied the nitration of benzene in nitro~benzene. but they did not ubtain reaction velocity constants. Besides these communications in the older literature there have. so faro only appeared publications on this subject by MARTINSEN 4). KLEMENC 5). and WIBAUT 6). .

MARTINSEN has measured the nitration velocity of some nitro~ and dinitro~derivatives of benzene in strong sulphuric acid. and also of phenol and p~cresol in water. The values for k (in sulphuric acid). calculated according to the bimolecular reaction equation RH + HN03 .... R . N02 + H 20 present a very satisfactory constancy at 25° and 46°; measured at 0°. however. they sensibly diminish in the course of the experiment. especially in nitro~benzene. MARTINSEN. therefore. considers the possibility that the reaction af ter all does not quite perfectly ag ree with the bimolecular scheme. He has further examined the influence of nitrous acid on the reaction velocity. It appeared to be equal to zero here.

The nitration of phenol and p~cresol in water has a complicated course; by~products are formed. and the experiment yields no constant values for k. points, on the other hand. to a distinct auto~catalytic course of the reaction. This is to be attributed to the nitrous acid that is formed in the liquid (proved colorimetrically according to ERDMAI'.iN 7). as the addition of only slight quantities of nitrite al ready considerably increases

1) This subject was given me by Prof. WIBAUT to treat as a thesis for the doctorate; I started the measurements communicated here, when I was employed as an assistant at the physiological laboratory of Amsterdam; when this had come to an end Prof. VAN

RIJNBERK gave me an opportunity to continue my investigation; I gladly express my indebtedness to him for this kindness.

2) Z. Phys. Ch. 2, 676 [1888]. 3) Dissert. Heidelberg [1899] . • ) Z. Phys. Ch. 50, 385 [1905]; 59, 605 [1907]. S) Monatsh. f. Ch. 35, 151 [1914]; 39, 151 [1918J; Z. f. anorg. u. allgem. Ch. Hl,

231 [192i]. 6) Rec. 3., 211 [1915]. 7) B 33, 210 [1900].

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the velocity of reaction. If the HN02 is, however, removed by the addition of urea, no nitration takes pI ace. The molarity of the urea was 0.05, that of the HN03 and the C6H sOH 0.6 resp. 0.075.

KLEMENC has made measurements on phenols, dissolved in ether. The process of the reaction appeared, however, to be very complicated in these very readily oxidizable benzene derivatives; henc~, with a few exceptions, the author has not succeeded in obtaining nitration constants that were in any way satisfactory. In his latest treatise on this subject 1) he assumes that the nitration is due to an acid H 2N 30 g ;

its character is, ho wever, still pretty hypotheticaI. as KLEMENC has to admit himself.

The two principal facts established from this investigation are: 1. a solution of HN03 in ether which i~ perfectly free from HN02 or N02,

does not nitrate; 2. if the nitric acid is present in excess, no nitration takes place or only a slight one.

WIBAUT has measured the velocity of nitration of benzene, toluene, chloro- and bromo-benzene in anhydrous acetic acid; for toluene at 0°, for the other th ree substances at 25°. With a single exception the concentration of the HN03 was always ab out 0.75, that of the benzene derivative about 0.5. The calculation of the k's was made according to the bimolecular reaction equation. As appears from the tab les referring to this 2), the values of the k's in benzene and in chloro-benzene are pretty constant; the departures from the mean are, however, sometimes very great indeed (in benzene e.g. up to 28 0/ 0), but this is attributed by the author to errors of . observation. The means of the parallel experiments inter se are in much better agreement. In bromo-benzene the k's increase in the course of the reaction, in toluene they decrease. Por benzene and chloro-benzene at 25° the following values were found : k = 0.0025, resp. 0.0020.

11. The Author's Experiments.

These form a direct continuation of the above-mentioned investigation by WIBAUT. Most experiments have been made by me in anhydrous acetic acid as a solvent; I have, however, also performed a few measure­ments in glacial acetic acid.

The anhydrous acetic acid which I employed, was that of the firm of DE HAEN "doppelt gereinigt". The reactions to detect Cl- and S01-­were negative; from analyses the content appeared to be about 98 %,

Some experiments in which I used MERCK' s preparation "Zur Analyse, frei von höheren Homologen", yielded no different results; DE HAEN's preparation is, accordingly, sufficiently pure for my purpose. The benzene was a perfectly pure preparation with a meIting point of 5°.48. I obtained

1) Z. für anorg. u. allgem. Chem. Hl. 236 [192-4]. 2) Compare for this the original.

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the absolute nitric acid by distillation of .pure 65 % HN03 with double the quantity by weight of strong sulphuric acid in vacuum in an apparatus entirely made of glass. The content was as a rule ab out 99.5 0/ 0' Af ter the preparation it was immediately placed in the ice-box.

Analysis of the samples. This was made by the method also used by WIBAUT. which consists in a quantity of the reaction mixture pipetted off. flowing into an excess of potassium hydroxide of the spec. weight of .1.3; the KN03 formed was then determined according to Devarda. af ter removal of the nitro-benzene formed through shaking out with benzene. By means of a number of experiments with a known quantity of KN03 I convinced myself th at the accuracy of this method is sufficient.

The róZe of the HN02• Considering the study of the röle of the HN02

in these experiments of great importance. I have tried to determine its concentration both at the beginning and during the reaction. This was done colorimetrically according to GRIEs-ILOSVAY 1) in the samples of the reaction mixture diluted in' a suitable way 2). The nitric acid was rendered innoxious by the addition of solid Na-acetate.

Result of some experiments. Table 1 renders an experiment at 25° 3);

the determinations were all made in duplicate; urea was not added; only the initial concentration of the HN02 was measured. Leaving the first measurement out of account. it is seen that the k (calculated according

h b I I 2.3025 C'oC) II to t e imo ecu ar equation k = t(Co-C'o)log CoC' remains very we

constant over a conversion range of more than 50 010 (probably over a still greater range: the interruption of the measurements during the night prevented me from ascertaining this). Not until the end of the reaction do the k's increase greatly, a fact which I found in almost all my experiments; though the relative errors at the beg inning and at the end of the reaction are, of course, much greater than in the middle part, this cannot account for the fact that the deviations are always found only in a positive sense.

The value I find for k, is therefore somewhat lower than that found by WIBAUT (0.0025).

I will now proceed to communicate an experiment, also made at 25°, but with addition of ± 8 mg of urea. This quantity is so sm all that in an analytical respect, it cannot exert any disturbing influence. HN02-

1) Cf. LUNGE-BERL. Chem.-techn. Unters. Meth. I. 853 [1921]. 2) In this hesides HN02 also N02(N20 4) will occur, which. on being poured out into

water, yields HN02 + HN03 ; hence the HN02-concentration found colorimetrically will he somewhat smaller than the sum of the concentrations [HN02] + IN02] in the reaction liquid. This has. however. no inlluence on the order of magnitude. In the tables 1 and 3 it has been assumed that exclusively HN02 is present.

3) The oscillations of temperature did not amount to more than ± 00 .02.

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determinations have not yet been carried out in this kind of experiments (TabIe 2). As is seen. the experiment again yields constants. but now the value of k is very much smaller. The conclusion is obvious that the HN02 considerably accelerates the reaction catalytically.

TABLE 1 (Benzene. temp. = 25° C.l

0

56.4

131.5

224

451

574

1352

1590

Initial concentrations HN03 = 0.6682 C6H6 = 0.4385 Co - C'o = 0.2297 HN02 = 0.0011

No urea added

c k

0.6682 -

0.6191 (0.00329)

0.5905 236

0.5529 225

0.4762 229

0.4417 237

0.2742 ( 450)

0.2420 ( 700)

Mean 0.00232 After 574' 51.6 Ofo converted Greatest deviation from the mean 3.0 0/0

T ABLE 2 (Benzene. temp. = 25° C,)

Initlal concentrations HN03 = 0.6808 C~6= 0.4808 Co - Cia = 0.2000

About 8 mg. of urea added

c k

0 0.6671 -

109.9 0.6230 0.00140

186.4 0.6001 131

293 0.5610 144

410 0.5282 145

546 0.5017 139

654 0.4484 ( 179)

1551 0.2627 ( 347)

Mean O.OOliO

Af ter 546' 37.2 Ofo converted Greatest deviation from the mean 6.4 0/0

Reaction of the HN03 with the solvent. Up to now no attention had been given to possible action of the HN03 caused by the anhydrous acetic acid itself. In the literature same records are, indeed, found of this reaction. chiefly in connection with the preparation of tetranitro methane 2), but the number of equivalents of the HN03 is then always of the same order of magnitude as of the anhydride. WIBAUT supposed 3) that ac etyl nitrate is formed in the nitration experiments described in his publication. which would th en be the active agent in the nitration.

Attempts to isolate this substance from a diluted solution of HN03

in anhydrous acetic acid have, however, not succeeded. However this may be, it appears from Table 6 of his publication 4) that na nitration

I) In these experiments time is always expressed in minutes. 2) Cf. e.g. PICTET and GENEQUAND. B. 36. 2225 [1910]; CHATTAWAY, Soc. 97. 2099

[191OJ. 3) Loc. cito p. 245. i) Loc. cito p. 250.

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takes place with excess of gramme~molecules of chloro~benzene; nor have I found nitration in a similar experiment with benzene. In these cases the concentration of the HN03 was. ho wever. much smaller than usual. From this the conclusion had been drawn that the titer of the HN03• wh en there is no benzene derivative present. does not decrease. It has. however. appeared to me from experiments made on purpose to study this. that it is very certainly the case. This can. of course. not be owing to a convers ion into acetyl nitrate. as this would have no influence on the titer found. In Tables 3 and -4 some experiments are recorded from which th is decrease of titer appears.

It will be seen th at this is very considerable. It may be read from the figure (curves 11 and 111) that the concentration remains constant for ab out 180 minutes. af ter which it decreases linearly with the time during a great interval. This is particularly clearly to be seen in curve 111. This need not mean that this relation is also theoretically true; the course of the curve sooner suggests the existence of two points of inflection. and it may therefore represent a consecutive reaction; practically. however. the middle portion is straight.

TABLE 3. TABLE 4.

(Experiment without benzene. temp. = 25° C.) (Experiment without benzene. temp. = 25° C.)

No urea added 15 mg. of urea added

CHNO, C

0 0.7142 0.00030 0 0.7221

34.7 0.7157 53 0.7168

105.9 0 . 7137 179 0.7185

142.7 0.7139 376 0.7018

189.7 0.7015 0.0089 584 0.6339

270 0.6754 1554 0.3338

335 0.6462 1811 0.2901

414 0.6181 0.0157

716 0.5121

720 0.5056 1)

909 0.4485

998 0.4216

1083 0.3921

1) New experiment. started the preceding day in the evening.

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If it is now assumed that the reactions of the HN03• on one side with the benzene. and on the other side with the solvent. proceed

independently of each other. this gives rise to the two simultaneous differential equations

dx - =k(A -x- y) (B- x) dt

(1)

(2)

in wich A and Bare the initial concentrations of nitric acid. resp. benzene. Equation (2) integrated yields y = at. as for t = 0 Y is = o. Substitution of this in (1) yields:

dx dt = k (A - x - at) (B ~ x) (3)

This differential equation would then have to represent the course of the nitration. at least from about 3 hours af ter the beginning of the

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reaction. We now see from Table 3 that in the interval of 720-909 minutes the titer of the HN03 decreases by an amount of 0.0571 mol.; per minute th is is 0.000302 mol. Table 4 gives this same amount over a still greater interval of conversion. We may therefore put a = 0.000302 in equation (3). If in Table 1 we now con si der the interval of 224-574 min., the HN03 would have been diminished by 350 X 0.000302 = - 0.1057 mol. through the reaction with the solvent.

The total decrease being 0.1112 mol., only 0.0055 mol. remains for the reaction with the benzene, hence an impossibly small amount. It is, indeed, easy to choose such a range of conversion that the decrease of the benzene concentration would change into an increase. It therefore appears (and the satisfactory constancy of the k's calculated without the assumption of a by~reaction teaches th is too) that equation (3) cannot. render the facts. Two explanations seem possible to me, viz.:

1. The reaction products of the HN03 and the anhydride themselves nitrate too; a priori this is, however, not very probable. Of a nitration with tetranitromethane a record is, indeed, found in the literature 1), but there either the benzene derivate itself was of a basic nature, or a basic substance (pyridine) was added to it. It is also possible that intermediately nitro acetic acid is formed, or another intermediate product; whether this could nitrate here I dare not decide. At any rate it would be exceedingly accidental that the effect should be exactly as if no by~ reaction had taken place.

2. The presence of the benzene prevents the by~reaction; this would e.g. be conceivable when an addition compound was formed. In this case the reaction would, ho wever, proceed monomolecularly; I have, therefore, calculated constants in some experiments according to the

2.3025 B equation k = --- log -B (B = initial concentration benzene, hence

t -x also that of the addition compound on excess of HN03); sometimes the k's were then actually in somewhat better agreement, but in the experiments performed with the greatest care, as in those of Table 1 this was very certainly not the case. It should be taken into account however that there will be an equilibrium between addition compound and its components; the reaction scheme would th en become: A + B ;: A . B ~ C + D.

Now whichever explanation may be the true one, I thought it in any case desirabie to determine besides the nitric acid, also the nitro~ benzene formed, because I have th en a direct answer to the question whether the HN03 is only abstracted by the benzene or also at the same time by the solvent. This has succeeded in the following way: from the sample received in lye from the reaction mixture, the nitro benzene was shaken out by petroleum ether, the solvent was distilled

1) E. SCHMIDT und U. FISCHER. Tetranitromethan als Nitrierungsmittel. B. 53. 1529 [1920).

699

off at a fractionation column and the C 6H sN02 was dissolved in alcohol. In this it was determined titanometrically according to KNECHT and HIBBERT 1) in the modified arrangement. as indicated by KOL THOFF 2). By analyses of solutions of C 6H sN02 in anhydrous acetic acid of known content I have beforehand ascertained the practicality of this method. I will confine myself here to stating this. as I have the results. obtained in a real nitration experiment. not yet completely at my disposal.

Ionisation of the nitric acid. It is. of course. conceivable that the HN03 in the anhydrous acetic acid is partially ionized. In this case the k' s would have to be calculated from a more complicated equation. as the effective concentration of the HN03 is smaller than the concentration determined analytieally. However WALDEN 3) has shown that the ionisation of acids in solvents is independent of the law of NERNST~ THOMSON: thus e of strong sulphuric acid is almost equal to that of water and yet the HN03 in these solvent is not split up into H+ and N03 : for. if this were the case. the reaction velocity constants found by MARTINSEN could not have been in such good agreement.

Reaction between anhydrous acetic acid and nitric acid with strong sulphuric acid as a catalyst. Wheri once to a solution of HN03 in (CH3COhO of the usual streng th (about 0.75 n) I added a little strong H 2SO" with the intention of diminishing the NOrcontent in this way. I suddenly observed an almost explosive reaction. The temperature of the vehemently boiling liquid rose toabout 122°. On repetition of the expe~ riment I found that af ter dilution with water an oil separates. whieh appeared to be tetranitromethane: the aqueous layer rendered alkaline contained NH3 ; the nitrie acid must. therefore. have acted very strongly oxydizing. whieh also appears from the expulsion of nitrie vapours. I further found that previous addition of urea prevents the reaction. or at least greatly retards it. If no H 2SO" is added. solutions of the concen~ tration used can remain at room temperature for a long time without anything being perceived of a violent reaction. The catalytie influence of the H 2SO" in this reaction ") has. as far as I have been able to verify. not been described in the literature as yet.

I have also made experiments in whieh the reaction flask was placed in a large dish filled with cold water. whieh was vigorously stirred in order to lead off the reaction heat quickly. Por I hoped th at the HN03 -

at least partly - would react according to the equation HN03 + (CH3COhO ~ CH3CON03 + CH3COOH.

In this way the acetyl nitrate could thee be prepared without it being

1) E. KNECHT and E. HIBBERT. New Reduction Methods in Volumetrie Analysis. 2) I. M. KOLTHOFF. Rec. 45. 169 (1926). 3) P. WALDEN. Trans. FARADAY Soc. 6. 71 [1910). 4) Addition of 8-10 dr. per 50 cc. solution is sufficient.

700

necessary to make N 20 5 first. A fractional vacuum distillation to which I subjected the mixture. has not produced the desired result so faro If possible I will repeat this experiment.

Summary of the Results.

At this place it will be sufficient for me to summarize my other experiments shortly. At 18° good constants were also obtained. and this over a wide conversion range; in one experiment k was found on an average 0.00128 ; from kl8 and k25 the value 2.3 follows for the temperature coefficient between 18° and 25°; at 0° and at "Wo the "constants" were bad; the k' s decrease. resp. increase considerably here. - In an experiment in which the benzene was present in an excess of gramme molecules. no nitration took pI ace ; here the concentrations applied were. however. small; it is. therefore. desirabIe to repeat this experiment with e.g. [C6H 61 = 0.75 and [HN03] = 0.5. - In another experiment the HN03 was ab out 0.5 molar. the C 6H 6 about 0.33. Here. too. constants were obtained but. very remarkably. k250 was here only 0.00166 instead of 0.00232. as in the experiment of Table 1. By measurement of the initial concen~ trations of the HN02 it will have to be decided wh ether they are possibly in (linear) relation to the reaction constants. Intentional addition of HN02 must render k greater then. - In glacial acetic acid no nitration takes place (or at least an exceedingly slow one): Af ter 24 hours the titer of the nitric acid is still unchanged. Addition of anhydride to bind the water in the glacial acetic acid did not change this at all.

In the experiments which I still hope to make. I shall in particular devote my attention to the röle of the concentration of the nitric acid in connection with the initial concentration of the HN02• It is further my intention to carry out a few more measurements with acetyl nitrate. both in anhydrous acetic acid and in tetrachloro~carbon. This is possible. as there is now no water-abstracting agent necessary. - In st rong sulphuric acid the nitration proceeds very rapidly. in ether on the other hand. very slowly; it is. therefore. possible th at these substances. applied in suitable mixing proportion. form a solvent in which the nitration of benzene. chloro-benzene etc. proceeds with measurable velocity. Also in this direction it will perhaps be tried to arrive at some results.

Amsterdam. 23. 6. '28. Physiol. Lab. of the University.